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Chapter 7: Problem 98

The electron affinities, in \(\mathrm{kJ} / \mathrm{mol}\), for the group 11 and group 12 metals are as follows: (a) Why are the electron affinities of the group 12 elements greater than zero? (b) Why do the electron affinities of the group 11 elements become more negative as we move down the group? [Hint: Examine the trends in the electron affinities of other groups as we proceed down the periodic table.]

Short Answer

Expert verified
(a) Group 12 elements have positive electron affinities because the incoming electron occupies a higher energy level orbital that is further away from the nucleus, resulting in a poor shielding effect and less attraction between the electron and nucleus. (b) The electron affinities of group 11 elements become more negative as we move down the group due to an increase in atomic radius, which leads to weaker attraction between valence electrons and the nucleus, allowing the atom to more easily gain an electron.

Step by step solution

01

Define Electron Affinity

Electron affinity is the energy change that occurs when an electron is added to a gaseous atom, forming an anion. It's a measure of the attraction between the nucleus and the incoming electron. A negative electron affinity value means that the atom releases energy when it gains an electron, while a positive value means that energy must be supplied to the atom for it to gain an electron.
02

Group 12 Elements Electron Affinities

Group 12 elements have electron affinities greater than zero, meaning they require energy to gain an electron. This is because when an electron is added to the atom, it occupies the (n+1) s orbital. The (n+1) s orbital is further away from the nucleus and has a higher energy level, leading to a poorer shielding effect. This results in less attraction between the incoming electron and the nucleus, and thus a positive electron affinity.
03

Group 11 Elements Electron Affinities

As we move down the group 11 elements in the periodic table, the electron affinities become more negative. This indicates that the atom releases more energy as it gains an electron. The atomic size increases down the group due to the addition of more electron shells and increased shielding effect of the inner electrons. This increase in atomic radius leads to an increased distance between the nucleus and the electrons in the outermost energy level (valence electrons). Consequently, valence electrons experience weaker attraction to the nucleus, which allows them to be more easily gained or removed.
04

Summarizing the Reasons

(a) Group 12 elements have electron affinities greater than zero due to poor shielding effect of the incoming electron, which occupies a higher energy level orbital further away from the nucleus. (b) Group 11 elements have more negative electron affinities down the group due to an increase in atomic radius, resulting in weaker attraction between valence electrons and the nucleus, making it easier for the atom to gain an electron.

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Most popular questions from this chapter

(a) One of the alkali metals reacts with oxygen to form a solid white substance. When this substance is dissolved in water, the solution gives a positive test for hydrogen peroxide, \(\mathrm{H}_{2} \mathrm{O}_{2}\). When the solution is tested in a burner flame, a lilac-purple flame is produced. What is the likely identity of the metal? (b) Write a balanced chemical equation for the reaction of the white substance with water.

Identify each statement as true or false: (a) lonization energies are always endothermic. (b) Potassium has a larger first ionization energy than lithium. (c) The second ionization energy of the sodium atom is larger than the second ionization energy of the magnesium atom. (d) The third ionization energy is three times the first ionization energy of an atom.

Elemental barium reacts more violently with water than does elemental calcium. Which of the following best explains this difference in reactivity? (i) Calcium has greater metallic character than does barium. (ii) The electron affinity of calcium is smaller than that of barium. (iii) The first and second ionization energies of barium are less than those of calcium. (iv) The atomic radius of barium is smaller than that of calcium. (v) The ionic radius of the barium ion is larger than that of the calcium ion.

Write the electron configurations for the following ions, and determine which have noble-gas configurations: (a) \(\mathrm{Ti}^{2+},(\mathbf{b})\) (d) \(\mathrm{PO}^{2-}\), (f) \(\mathrm{V}^{3+}\) \(\mathrm{Br}^{-}\) (c) \(\mathrm{Mg}^{2+}\) (e) \(\mathrm{Pt}^{2+}\)

Consider the \(\mathrm{A}_{2} \mathrm{X}_{4}\) molecule depicted here, where \(\mathrm{A}\) and \(\mathrm{X}\) are elements. The \(A-A\) bond length in this molecule is \(d_{1}\), and the four \(\mathrm{A}-\mathrm{X}\) bond lengths are each \(d_{2}\). (a) In terms of \(d_{1}\) and \(d_{2},\) how could you define the bonding atomic radii of atoms A and X? (b) In terms of \(d_{1}\) and \(d_{2}\), what would you predict for the \(X-X\) bond length of an \(X_{2}\) molecule? [Section \(\left.7.3\right]\)
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